Important Compounds

Among the most commercially and analytically significant compounds of the d-block (transition) elements are potassium dichromate ( $\ce{K_2Cr_2O_7}$ ) and potassium permanganate ( $\ce{KMnO_4}$ ). These vivid, crystalline solids—orange and deep purple, respectively—owe their striking colours to d–d electronic transitions and charge-transfer processes within the chromium(VI) and manganese(VII) centres. Both compounds are powerful oxidising agents, a property that underpins their extensive use in analytical chemistry, industrial processes, and organic synthesis. Potassium dichromate is prepared industrially from chromite ore $\ce{FeCr_2O_4}$ through a sequence of roasting, leaching, and acidification steps. The chromium is ultimately stabilised in the hexavalent (+6) oxidation state as the dichromate ion $\ce{Cr_2O_7^{2-}}$ . Potassium permanganate, on the other hand, is obtained from pyrolusite ( $\ce{MnO_2}$ ) by alkaline fusion followed by electrolytic or chemical oxidation, bringing manganese to its highest common oxidation state of +7 in the permanganate ion $\ce{MnO_4^-}$ . Understanding the preparation routes helps explain the structural and electronic properties of each compound. In quantitative analysis, both compounds are primary or secondary standards in volumetric (titrimetric) analysis. Permanganate titrations are self-indicating due to the intense purple colour of $\ce{KMnO_4}$ versus the near-colourless $\ce{Mn^{2+}}$ product in acid solution, making them especially elegant. Dichromate titrations, while requiring an external indicator such as diphenylamine, offer greater stability of the standard solution. Mastery of these redox titrations—including half-equations, mole ratios, and equivalence-point calculations—is central to A-level and IB chemistry courses.

Key takeaways

Industrial Preparation of K₂Cr₂O₇
Structure and Properties of K₂Cr₂O₇
Industrial Preparation of KMnO₄
Structure and Properties of KMnO₄
Permanganate Titrations (Acidic Medium)
Dichromate Titrations

Industrial Preparation of K₂Cr₂O₇

Step 1 — Roasting chromite with sodium carbonate in excess air at ~1000 °C converts Cr(III) to soluble sodium chromate: $\ce{4FeCr_2O_4 + 8Na_2CO_3 + 7O_2 -> 8Na_2CrO_4 + 2Fe_2O_3 + 8CO_2}$ . Step 2 — The yellow chromate solution is acidified (H₂SO₄) to shift the equilibrium toward the orange dichromate: $\ce{2CrO_4^{2-} + 2H^+ <=> Cr_2O_7^{2-} + H_2O}$ . Step 3 — Addition of KCl or K₂SO₄ precipitates less-soluble $\ce{K_2Cr_2O_7}$ , which is filtered and recrystallised to give the pure orange solid (mp 398 °C).

Structure and Properties of K₂Cr₂O₇

The dichromate ion consists of two corner-sharing $\ce{CrO_4}$ tetrahedra linked through a bridging oxygen, giving a Cr–O–Cr bond angle of ~126°. Each chromium is in the +6 state (d⁰ configuration). The intense orange colour arises from ligand-to-metal charge transfer (LMCT) rather than d–d transitions (d⁰ systems have no d–d transitions). In acid solution, $\ce{K_2Cr_2O_7}$ is a strong oxidising agent: $\ce{Cr_2O_7^{2-} + 14H^+ + 6e^- -> 2Cr^{3+} + 7H_2O}$ , $E^\circ = +1.33\text{ V}$ . It oxidises Fe²⁺ → Fe³⁺, I⁻ → I₂, and is used in the breathalyser reaction (ethanol → acetic acid).

Industrial Preparation of KMnO₄

Step 1 — Pyrolusite ( $\ce{MnO_2}$ ) is fused with $\ce{KOH}$ and an oxidant ( $\ce{KNO_3}$ or air) at ~200 °C to form green potassium manganate(VI): $\ce{2MnO_2 + 4KOH + O_2 -> 2K_2MnO_4 + 2H_2O}$ . Step 2 — The manganate is oxidised to permanganate either electrolytically (anode: $\ce{MnO_4^{2-} -> MnO_4^- + e^-}$ ) or by disproportionation in dilute acid: $\ce{3MnO_4^{2-} + 4H^+ -> 2MnO_4^- + MnO_2 + 2H_2O}$ . Step 3 — The deep purple $\ce{KMnO_4}$ crystallises on cooling and is purified by recrystallisation.

Structure and Properties of KMnO₄

The permanganate ion $\ce{MnO_4^-}$ is tetrahedral with Mn in the +7 oxidation state (d⁰). Its vivid purple colour results from intense LMCT bands (O→Mn) in the visible region. In acidic solution the standard reduction potential is very high: $\ce{MnO_4^- + 8H^+ + 5e^- -> Mn^{2+} + 4H_2O}$ , $E^\circ = +1.51\text{ V}$ . In neutral/alkaline solution the product is $\ce{MnO_2}$ (brown precipitate, Mn = +4), and in strongly alkaline solution it forms $\ce{MnO_4^{2-}}$ (Mn = +6). This pH-dependence of products is a key examination topic.

Permanganate Titrations (Acidic Medium)

$\ce{KMnO_4}$ is used as an oxidising titrant in H₂SO₄ solution. Common analytes include: (a) Iron(II): $\ce{MnO_4^- + 5Fe^{2+} + 8H^+ -> Mn^{2+} + 5Fe^{3+} + 4H_2O}$ ; (b) Oxalate: $\ce{2MnO_4^- + 5C_2O_4^{2-} + 16H^+ -> 2Mn^{2+} + 10CO_2 + 8H_2O}$ ; (c) Hydrogen peroxide: $\ce{2MnO_4^- + 5H_2O_2 + 6H^+ -> 2Mn^{2+} + 5O_2 + 8H_2O}$ . The endpoint is signalled by the first permanent pale pink colour from excess $\ce{KMnO_4}$ . No separate indicator is needed — the titration is self-indicating. Mole ratios (1:5 for Mn:Fe; 2:5 for Mn:oxalate) are critical for calculations.

Dichromate Titrations

$\ce{K_2Cr_2O_7}$ is a primary standard (stable, pure, non-hygroscopic, known formula weight = 294.18 g mol⁻¹). A typical reaction in acid is with iron(II): $\ce{Cr_2O_7^{2-} + 6Fe^{2+} + 14H^+ -> 2Cr^{3+} + 6Fe^{3+} + 7H_2O}$ . The mole ratio is 1:6 (Cr₂O₇²⁻ : Fe²⁺). The endpoint is detected with diphenylamine indicator (colourless → violet) or by potentiometric methods. Because $\ce{K_2Cr_2O_7}$ solutions are indefinitely stable, they are preferred over $\ce{KMnO_4}$ (which slowly decomposes in light and oxidises water) when a primary standard is needed.

The Key Equation

\\ce{MnO_4^- + 5Fe^{2+} + 8H^+ -> Mn^{2+} + 5Fe^{3+} + 4H_2O}

\ce{MnO_4^-} — Permanganate ion (mol): The oxidising agent; Mn is reduced from +7 to +2, gaining 5 electrons per ion.
\ce{Fe^{2+}} — Iron(II) ion (mol): The reducing agent (analyte); Fe is oxidised from +2 to +3, losing 1 electron per ion.
\ce{H^+} — Hydrogen ion (mol): Provided by dilute H₂SO₄; essential for the reaction — HCl is avoided as Cl⁻ can be oxidised by KMnO₄.
\ce{Mn^{2+}} — Manganese(II) ion (mol): Nearly colourless product; its formation causes the purple colour of KMnO₄ to discharge during the titration.

Worked Example

Problem. 25.00 cm³ of an acidified iron(II) sulfate solution required 22.50 cm³ of 0.0200 mol dm⁻³ KMnO₄ for complete oxidation. Calculate the concentration of Fe²⁺ in the solution.

c(\ce{Fe^{2+}}) = 0.0900 \text{ mol dm}^{-3}

Answer. The concentration of Fe²⁺ is 0.0900 mol dm⁻³.

Figure

Diagram

Origins · 1 minJohann Rudolf Glauber and how this was discovered

Johann Rudolf Glauber

1604–1670 · German

Glauber was one of the earliest chemists to systematically investigate the preparation and properties of inorganic salts and oxidising mineral compounds. His work on the preparation of various potassium and sodium salts laid foundational principles for understanding salt chemistry, directly anticipating the industrial preparation sequences later used for compounds like K₂Cr₂O₇ and KMnO₄. His treatise 'Furni Novi Philosophici' (1648) described furnace reactions—analogous to the roasting and fusion steps used today. Glauber discovered sodium sulfate (Na₂SO₄·10H₂O) which he called 'sal mirabile Glauberi' (Glauber's salt), believing it had miraculous medicinal properties—he used it as a universal remedy and reportedly made a fortune selling it.

Variable Oxidation States in Transition Metals — K₂Cr₂O₇ (Cr: +6) and KMnO₄ (Mn: +7) represent high oxidation states. Understanding why d-block elements exhibit variable oxidation states—due to comparable energies of (n−1)d and ns electrons—is essential context.
Redox Half-Equations and Standard Electrode Potentials — The oxidising power of both compounds is quantified by E° values. Writing and balancing half-equations is the core skill for predicting and calculating the outcomes of permanganate and dichromate titrations.
Volumetric Analysis and Titrations — Permanganate and dichromate titrations are classic examples of redox volumetric analysis. Understanding burette technique, equivalence points, and mole ratio calculations is directly required.
Colour in Transition Metal Complexes — The orange colour of Cr₂O₇²⁻ and purple of MnO₄⁻ both arise from LMCT rather than d–d transitions because both metal centres are d⁰. This connects to the broader theory of electronic transitions in coordination compounds.